Chemistry

Key unit competency

Interpret simple mass spectra and use them to calculate the relative atomic mass (R.A.M) of different elements.

Learning objectives

By the end of this unit, I will be able to:

•Outline the discovery of the sub-atomic particles.

•Compare the properties of sub-atomic particles.

•Explain what is an isotope of an element.

•Assess the relationship between the number of protons and the number of electrons.

•Calculate the mass number knowing the number of protons and the number of neutrons.

•Understand the meaning of relative atomic mass and relative abundances

•Calculate the relative atomic mass of an element, given isotopic masses and abundances.

•Draw and label the mass spectrometer.

•Explain the fundamental processes occurring in the functioning of a mass spectrometer.

• Interpret different mass spectra.

•State the uses of the mass spectrometer.

•Calculate the relative atomic mass of an element, from a mass spectrum.

Each country has its own culture (language, traditions and norms, attitudes and values, etc.). Our culture defines our identity which is unique to each Rwandan citizen and differentiates us from foreigners; if one element of our culture is rejected or disappears, we become a different Rwandan people. When we introduce foreign cultures to replace ours, we can lose our identity. However, some of our cultural elements such as language can be shared with others to build the social relationship.

Similarly, in the atom, the number of protons within the nucleus defines the atomic number, which is unique to each chemical element; the atomic number or the number of protons of an atom defines its identity. If a proton is added or removed from an element, it becomes a different element. Electrons around the nucleus can be lost, gained, or shared to create bonds with other atoms in chemical reactions to produce useful substances, but this does not change the identity of the elements involved.

1.1. Outline of the discovery of the atom constituents and their properties

Activity 1.1

1. Regardless of some exceptions, all atoms are composed of the same com-ponents. True or False? If this statement is true why do different atoms have different chemical properties?

2. The contributions of Joseph John Thomson and Ernest Rutherford led the way to today’s understanding of the structure of the atom. What were their contributions?

3. Explain the modern view of the structure of the atom?

4. Using your knowledge about atom, what is the role each particle plays in an atom?

1.1.1. Constituents of atoms and their properties

Atoms are the basic units of elements and compounds. In ordinary chemical reactions, atoms retain their identity. An atom is the smallest identifiable unit of an element. There are about 91 different naturally occurring elements. In addition, scientists have succeeded in making over 20 synthetic elements (elements not found in nature but produced in Laboratories of Reasearch Centers).

An element is defined as a substance that cannot be broken down by ordinary chemical methods in simpler substances. Some examples of elements include hydrogen (H), helium (He), potassium (K), carbon (C), and mercury (Hg). In an element, all atoms have the same number of protons or electrons although the number of neutrons canvary. A substance made of only one type of atom is called also element or elemental substance, for example: hydrogen (H2), chlorine (Cl2), sodium (Na). Elements are the basic building blocks of more complex matter.

A compound is a matter or substance formed by the combination of two or more different elements in fixed ratios. Consider, Hydrogen peroxide (H2O2) is a compound composed of two elements, hydrogen and oxygen, in a fixed ratio (2:2).

During the early twentieth century, scientists discovered that atoms can be divided into more basic particles. Their findings made it clear that atoms contain a central portion called the nucleus. The nucleus contains protons and neutrons. Protons are positively charged, and neutrons are neutral. Whirling about the nucleus are particles called electrons which are negatively charged. The relative masses and charges of the three fundamental particles are shown in Table 1.1

The mass of an electron is very small compared with the mass of either a proton or a neutron.

The charge on a proton is equal in magnitude, but opposite in sign, to the charge on an electron.

1.1.2. Discovery of the atom constituents.

The oldest description of matter in science was advanced by the Greek philosopher Democritus in 400 BC.He suggested that matter can be divided into small particles up to an ultimate particle that cannot any more be divided, and called that particle atom. Atoms came from the Greek word atomos meaning indivisible.

The work of Dalton and other scientists such as Avogadro, etc., contributed more so that chemistry was beginning to be understood. They proposed new concept of atom, and from that moment scientists started to think about the nature of the atom. What are the constituents of an atom, and what are the features that make atoms of the various elements to differ?

In 1808 Dalton published A New System of Chemical Philosophy, in which he presented his theory of atoms:

a) Dalton’s Atomic Theory

1. Each element is made up of tiny particles called atoms.

2. The atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways.

3. Chemical compounds are formed when atoms of different elements combine with each other. A given compound always has the same relative numbers and types of atoms.

4. Chemical reactions involve reorganization of the atoms—changes in the way they are bound together. The atoms themselves are not changed in a chemical reaction.

b)Discovery of Electrons and Thomson’s Atomic Model

In 1897 J. J. Thomson (1856–1940) and other scientists conducted several experiments, and found that atoms are divisible. They conducted experiments with gas discharge tubes. A gas discharge tube is shown in Figure 1.2.

The gas discharge tube is an evacuated glass tube and has two electrodes, a cathode (negative electrode) and an anode (positive electrode). The electrodes are connected to a high voltage source. Inside the tube, an electric discharge occurs between the electrodes.

The discharge or ‘rays’ originate from the cathode and move toward the anode, and hence are called cathode rays. Using luminescent techniques, the cathode rays are made visible and it was found that these rays are deflected away from negatively charged plates. The scientist J. J. Thomson concluded that the cathode ray consists of negatively charged particles, and he called them electrons.

Thomson postulated that an atom consisted of a diffuse cloud of positive charge with the negative electrons embedded randomly in it. This model, shown in Figure 1.3, is often called the plum pudding model because the electrons are like raisins dispersed in a pudding (the positive charge cloud), as in plum pudding.

In 1909 Robert Millikan (1868–1953) conducted the famous charged oil drop experiment and came to several conclusions: He found the magnitude of the charge of an electron equal to -1.602 x 10-19c. From the charge-to-mass ratio(e/m) determined by Thomson, the mass of an electron was also calculated.

c)Discovery of Protons and Rutherford’s Atomic Model

In 1886 Eugene Goldstein (1850–1930) observed that a cathode-ray tube also generates a stream of positively charged particles that move towards the cathode. These were called canal rays because they were observed occasionally to pass through a channel, or “canal,” drilled in the negative electrode (Figure 1.4). These positive rays, or positive ions, are created when the gaseous atoms in the tube lose electrons. Positive ions are formed by the process

Different elements give positive ions with different e/m ratios. The regularity of the e/m values for different ions led to the idea that there is a subatomic particle with one unit of positive charge, called the proton. The proton is a fundamental particle with a charge equal in magnitude but opposite in sign to the charge on the electron. Its mass is almost 1836 times that of the electron.

The proton was observed by Ernest Rutherford and James Chadwick in 1919 as a particle that is emitted by bombardment of certain atoms with α-particles.

Rutherford reasoned that if Thomson’s model were accurate, the massive α-particles should crash through the thin foil like cannonballs through gauze, as shown in Figure 1.6(a). He expected α-particles to travel through the foil with, at the most, very minor deflections in their paths. The results of the experiment were very different from those Rutherford anticipated. Although most of the α- particles passed straight through, many of the particles were deflected at large angles, as shown in Figure 1.6(b), and some were reflected, never hitting the detector. This outcome was a great surprise to Rutherford. Rutherford knew from these results that the plum pudding model for the atom could not be correct. The large deflections of the α-particles could be caused only by a center of concentrated positive charge that contains most of the atom’s mass, as illustrated in Figure 1.6(b). Most of the α-particles pass directly through the foil because the atom is mostly open space. The deflected α-particles are those that had a “close encounter” with the massive positive center of the atom, and the few reflected α-particles are those that made a “direct hit” on the much more

massive positive center.

In Rutherford’s mind these results could be explained only in terms of a nuclear atom—an atom with a dense center of positive charge (the nucleus) with electrons moving around the nucleus at a distance that is large relative to the nuclear radius.

d)Discovery of Neutrons

In spite of the success of Rutherford and his co-workers in explaining atomic structure, one major problem remained unsolved.

If the hydrogen contains one proton and the helium atom contains two protons, the relative atomic mass of helium should be twice that of hydrogen. However, the relative atomic mass of helium is four and not two.

This question was answered by the discovery of James Chadwick, English physicist who showed the origin of the extra mass of helium by bombarding a beryllium foil with alpha particles. (See figure 1.7).

In the presence of beryllium, the alpha particles are not detected; but they displace

uncharged particles from the nuclei of beryllium atoms. These uncharged particles cannot be detected by a charged counter of particles.

However, those uncharged particles can displace positively charged particles from another substance. They were called neutrons.The mass of the neutron is slightly greater than that of proton.

Figure 1.8 shows the location of the elementary particles (protons, neutrons, and electrons) in an atom. There are other subatomic particles, but the electron, the proton, and the neutron are the three fundamental components of the atom that are important in chemistry.

Atoms consist of very small, very dense nuclei surrounded by clouds of electrons at relatively great distances from the nuclei. All nuclei contain protons; nuclei of all atoms except the common form of hydrogen also contain neutrons.

Checking up 1.1

1. In an experiment, it was found that the total charge on an oil drop was 5.93 × 10-18 C. How many negative charges does the drop contain?

2. All atoms of the elements contain three fundamental particles. True or false? Give an example to support your answer.

3. Compare the atom constituents

a. in terms of their relative masses

b. in terms of their relative charges

4. Using the periodic table as a guide, specify the number of protons and electrons in a neutral atom of each of these elements.a. carbon (C) b. calcium (Ca) c. chlorine (Cl) d. chromium (Cr)

1.2. Concept of atomic number, mass number, and isotopic mass

Activity 1.2

The diagram below shows a representation of sodium isotopes . Observe it and answer to the questions below

1. Compare the two sodium isotopes in the figures above.

2. From your observation, how do you define the isotopes of an element?

3. How is the mass number, A, determined?

4. What information is provided by the atomic number, Z?

5. What is the relationship between the number of protons and the number of electrons in an atom?

6. Where are the electrons, protons, and neutrons located in an atom?

7. Why is the mass of an atom concentrated in the center?

8. Sodium-24 and sodium-23 react similarly with other substances. Explain the statement

9. Say which one(s) of the following statements is(are) correct and which one(s) is(are) wrong: (i) isotopes differ in their number of electrons, (ii) isotopes differ in their mass numbers, (iii) isotopes differ in their number of protons, (iv) isotopes differ by their number of neutrons, (v) all the statements are wrong.

The atomic number denotes the number of protons in an atom’s nucleus. The mass number denotes the total number of protons and neutrons. Protons and neutrons are often called nucleons. By convention, the atomic number is usually written to the left subscript of the elemental notation, and the mass number to the left superscript of the elemental notation as represented by the example below, where X represents any elemental symbol.

Some atoms of the same of element have the same atomic number, but different mass numbers. This means a different number of neutrons. Such atoms are called isotopes of the element.

Isotopes are atoms of the same element with different masses; they are atoms containing the same number of protons but different numbers of neutrons.

In a given atom, the number of protons, also called “atomic number” are equal to the number of electrons because the atom is electrically neutral. The sum of the number of protons and neutrons in an atom gives the mass number of that atom.

Mass number = number of protons + number of neutrons

                       = atomic number + neutron number

Checking up 1.2

1. How do you call the members of each of the following pairs? Explain.

2. Write, using the periodic table, the correct symbols to identify an atom that contains

a. 4 protons, 4 electrons, and 5 neutrons;

b. 23 protons, 23 electrons, and 28 neutrons;

c. 54 protons, 54 electrons, and 70 neutrons; and

d. 31 protons, 31 electrons, and 38 neutrons.

3. Use the list of the words given below to fill in the blank spaces. Each word will be used once.

Atomic number, Mass number, protons, Electrons, Isotopes, neutron

a) The atomic number tells you how many.................................. and ............................................................. are in an atom.......................................................is the number written as subscript on the left of the atomic symbol.

b) The total number of protons and neutrons in an atom is called the .......................................................................

c) Atoms with the same number of protons but different number of neu-trons are called .................................................

d) The subatomic particle that has no charge is called a .......................................................

1.3. Calculation of relative atomic mass of elements with isotopes

Activity 1.3

1. Argon has three naturally occurring isotopes: argon-36, argon-38, and argon-40. Based on argon’s reported relative atomic mass from the periodic table, which isotope do you think is the most abundant in nature? Explain.

2. Calculate the average atomic mass of an element with two naturally occurring isotopes: 85X (72.15%, 84.9118 amu) and 87X (27.85%, 86.9092 amu). Identify this element?

3. Boron has two naturally occurring isotopes. Find the percent abundances of 10B and 11B given the isotopic mass of 10B = 10.0129 amu and the isotopic mass of 11B = 11.0093 amu.

Relative atomic mass, symbolized as R.A.M or Ar, is defined as the mass of one atom of an element relative to 1/12 of the mass of an atom of carbon-12, which has a mass of 12.00 atomic mass units. The relative atomic mass, also known as the atomic weight or average atomic weight, is the average of the atomic masses of all of the element’s isotopes.

Relative isotopic mass is like relative atomic mass in that it deals with individual isotopes. The difference is that we are dealing with different forms of the same element but with different masses.

Thus, the different isotopic masses of the same elements and the percentage abundance of each isotope of an element must be known in order to accurately calculate the relative atomic mass of an element.

Notice: Remember that mass number is not the same as the relative atomic mass or isotopic mass! The mass number is the number of protons + neutrons; while relative atomic mass (or isotopic mass) is the mass if you were to somehow weigh it on a balance.

Let A1, A2, A3,..., An be an abundance of n isotopes of the same chemical element with atomic mass M1, M2, M3,..., Mn respectively, the relative atomic mass(R.A.M) is given by the following equation:

Example 1: Oxygen contains three isotopes 16O, 17O, and 18O. Their respective relative abundances are 99.76%, 0.04%, and 0.20%. Calculate the relative atomic mass of oxygen.

Solution:

Relative isotopic mass of 16O is 16 and its relative abundance is 99.76%;

Relative isotopic mass of 17O is 17, abundance 0.04%;

Relative isotopic mass of 18O is 18, abundance 0.20%.

By applying the same formula, the relative abundance of the isotopes may be calculated knowing the relative atomic mass of the element and the atomic masses of the respective isotopes.

Example 2: Chlorine contains two isotopes 35Cl and 37Cl, what is the relative abundance of each isotope in a sample of chlorine if its relative atomic mass is 35.5?

Solution:

Checking up 1.3

1. Three isotopes of magnesium occur in nature. Their abundances and masses, determined by mass spectrometry, are listed in the following table. Use this information to calculate the relative atomic mass of magnesium.

2.The atomic weight of gallium is 69.72 amu. The masses of the naturally occurring Isotopes are 68.9257 amu for 69Ga and 70.9249 amu for 71Ga. Calculate the percent abundance of each isotope

1.4. Description and functioning of the mass spectrometer

Activity 1.4:

By using the information in this book and other sources, attempt to answer the following questions

1. Define the mass spectrometer and state the main stages of its functioning.

2. Here below is given the block diagram of the mass spectrometer. Identify the unidentified marked component.

a)Inlet system

b)Ionisation chamber

c)Vacuumsystem

d) Ion transducer

3. Inlet system is also known as which of the following?

a)Initial system

b)Sample reservoir

c) Sample handling system

d) Element injection system

The mass spectrometer is an instrument that separates positive gaseous atoms and molecules according to their mass-charge ratio and that records the resulting mass spectrum.

In the mass spectrometer, atoms and molecules are converted into ions. The ions are separated as a result of the deflection which occurs in a magnetic and electric field.

The basic components of a mass spectrometer are: vaporisation chamber (to produce gaseous atoms or molecules), ionization chamber (to produce positive ions), accelerating chamber (to accelerate the positive ions to a high and constant velocity), magnetic field (to separate positive ions of different m/z ratios), detector (to detect the number and m/z ratio of the positive ions) and the recorder (to plot the mass spectrum of the sample).

A mass spectrometer works in five main stages, namely vaporization, ionization, acceleration, deflection, and detection to produce the mass spectrum.

Stage 1: Vaporization

At the beginning the test sample is heated until it becomes vapour and is introduced as a vapour into the ionization chamber. When a sample is a solid with low vapour pressure, it can directly be introduced into the ionization chamber.

Stage 2: Ionisation

The vaporized sample passes into the ionization chamber (with a positive voltage of about 10,000 volts). The electrically heated metal coil gives off electrons which are attracted to the electron trap which is a positively charged plate.

The particles in the sample (atoms or molecules) are therefore bombarded with a stream of electrons (electrons gun), and some of the collisions are energetic enough to knock one or more electrons out of the sample particles to make positive ions. Mass spectrometers always work with positive ions.

Most of the positive ions formed will carry a charge of +1 because it is much more difficult to remove further electrons from an already positive ion.

Most of the sample molecules are not ionized at all but are continuously drawn off by vacuum pumps which are connected to the ionization chamber (figure 1.9). Some of the molecules are converted to negative ions through the absorption of electrons.

The repeller plate absorbs these negative ions. A small proportion of the positive ions which are formed may have a charge greater than one (a loss of more than one electron). These are accelerated in the same way as the singly charged positive ions.

Stage 3: Acceleration

The positive ions are accelerated by an electric field so that they move rapidly through the machine at high and constant velocity.

Stage 4: Deflection

The ions are then deflected by a magnetic field according to their masses and charges ratio. Different ions are deflected by the magnetic field at different extents. The extent to which the beam of ions is deflected depends on four factors:

1. The magnitude of the accelerating voltage (electric field strength). Higher voltages result in beams of more rapidly moving particles to be deflected less than the beams of the more slowly moving particles produced by lower voltages.

2. Magnetic field strength. Stronger fields deflect a given beam more than weaker fields.

3. Masses of the particles. Because of their inertia, heavier particles are deflected less than lighter particles that carry the same charge.4. Charges on the particles. Particles with higher charges interact more strongly with magnetic fields and are thus deflected more than particles of equal mass with smaller charges

The two last factors (mass of the ion and charge on the ion) are combined into the mass/charge ratio. Mass/charge ratio is given the symbol m/z (or sometimes m/e).For example, if an ion had a mass of 28 and a charge of 1+, its mass/charge ratio would be 28. An ion with a mass of 56 and a charge of 2+ would also have a mass/charge ratio of 28.

In the figure 1.11 above, ion stream A is most deflected: it will contain ions with the smallest mass/charge ratio. Ion stream C is the least deflected: it contains ions with the greatest mass/charge ratio. Assuming 1+ ions, stream A has the lightest ions, stream B the next lightest and stream C the heaviest. Lighter ions are going to be more deflected than heavy ones.

Stage 5: Detection

The beam of ions passing through the machine is detected electrically. As they pass out of the magnetic field, ions are detected by an ion detector which records the position of the ions on the screen and the number of ions that hit the screen at each position. These two pieces of information are used to produce a mass spectrum for the sample.

A flow of electrons in the wire is detected as an electric current which can be amplified and recorded. The more ions arriving, the greater the current

Detecting the other ions

How might the other ions be detected (those in streams A and C which have been lost in the machine)?

Remember that stream A was most deflected. To bring them on to the detector, you would need to deflect them less by using a smaller magnetic field.

To bring those with a larger m/z value (the heavier ions if the charge is +1) to the detector you would have to deflect them more by using a larger magnetic field.If you vary the magnetic field, you can bring each ion stream in turn on the detector to produce a current which is proportional to the number of ions arriving. The mass of each ion being detected is related to the size of the magnetic field used to bring it on to the detector. The machine can be calibrated to record current (which is a measure of the number of ions) against m/z directly. The mass is measured on the 12C scale.

Note: The 12C scale is a scale on which the 12C isotope weighs exactly 12 units..

Recorder

The detector of a typical instrument consists of a counter which produces a current that is proportional to the number of ions which strike it. Through the use of electron multiplier circuits, this current can be measured so accurately that the current caused by just one ion striking the detector can be measured. The signal from the detector is fed to a recorder, which produces the mass spectrum. In modern instruments, the output of the detector is fed through an interface to a computer. The computer can store the data, provide the output in both tabular and graphic forms, and compare the data to standard spectra, which are contained in spectra libraries also stored in the computer.

This is an example of an appearance of a mass spectrum of unknown element that has 2 isotopes.

Checking up 1.4

1. Use the list of the words given below to fill in the blank spaces. Each word will be used once.

Vaporization chamber, mass spectrum, velocity, ionization, deflection, detector, acceleration

A sample of the element is placed in the _________ chamber where it is converted into gaseous atoms. The gaseous atoms are ionized by bombardment of high energy electrons emitted by a hot cathode to become positive ions (in practice, the voltage in the ________chamber is set in such a way that only one electron is removed from each atom). The positive ions (with different masses) are then going faster to a high and constant _________by two negatively charged plates: the process is called_________. The positive ions are then deviatedby the magnet field. This process is called ____________ (ions with smaller mass will be deflected more than the heavier ones). These ions are then detected by the ion _________. The information is fed into a computer which prints out the________ of the element.

2. The correct order for the basic features of a mass spectrometer is...

a. acceleration, deflection, detection, ionization

b. ionisation, acceleration, deflection, detection

c. acceleration, ionisation, deflection, detection

d. acceleration, deflection, ionisation, detection

3. Which one of the following statements about ionisation in a mass spectrometer is incorrect?

a. gaseous atoms are ionised by bombarding them with high energy electrons

b. atoms are ionised so they can be accelerated

c. atoms are ionised so they can be deflected

d. it doesn’t matter how much energy you use to ionise the atoms

4. The path of ions after deflection depends on...

a. only the mass of the ion

b. only the charge on the ionc. both the charge and the mass of the iond. neither the charge nor the mass of the ion

5. Which of the following species will be deflected to the greatest extent?

a. 37Na+

b. 35Na+

c. 37Na

d. 35Na2+

6. Which of the following separates the ions according to their mass-to-charge?

a) Ion source

b) Detector

c) Magnetic sector

d) Electric sector

1.5. Interpretation of mass spectra.

Activity 1.5

The mass spectrum of zirconium looks like this:

a) What does m/z mean?

b) Explain as fully as possible what the mass spectrum shows about zirconium.

(I am not expecting you to read actual values from the relative abundance axis.)

c) The spectrum shows lines for 1+ ions. If there were also peaks for 2+ ions, where would you expect to find them, and what would you predict about their heights relative to the 1+ peaks?

Example 1: The mass spectrum of boron

The mass spectrum of boron may be used to know the number of boron isotopes and their relative abundances

The two peaks in the mass spectrum shows that there are 2 isotopes of boron with relative isotopic masses of 10 and 11 on the 12C scale

The mass spectrum of an element shows how you can find out the masses and relative abundances of the various isotopes of the element and use that information to calculate the relative atomic mass of the element.

The relative size of the peaks gives you a direct measure of the relative abundances of the isotopes. The tallest peak is often given an arbitrary height of 100 but you may find all sorts of other scales used; it doesn’t matter. You can find the relative abundances by measuring the lines on the stick diagram.

In this case, the two isotopes (with their relative abundances) are:

Example 2: The mass spectrum for zirconium

The mass spectrum of zirconium may be used to know the number of zirconium isotopes and their relative abundances

The 5 peaks in the mass spectrum show that there are 5 isotopes of zirconium with relative isotopic masses of 90, 91, 92, 94 and 96 on the 12C scale.

This time, the relative abundances are given as percentages. Again you can find these relative abundances by measuring the lines on the stick diagram. In this case, the 5 isotopes (with their relative percentage abundances) are:

Example 3: The mass spectrum of chlorine

Chlorine is taken as typical of elements with more than one atom per molecule. Chlorine has two isotopes, 35Cl and 37Cl, in the approximate ratio of 3 atoms of 35Cl to 1 atom of 37Cl. You might suppose that the mass spectrum would look like this:

But it is not true. The problem is that chlorine consists of molecules, not individual atoms. When chlorine is passed into the ionization chamber, an electron is knocked off the molecule to give a molecular ion, Cl2+. These ions won’t be particularly stable, and some will fall apart to give a chlorine atom and a Cl+ ion. The term for this is fragmentation.

If the Cl atom formed isn’t then ionized in the ionization chamber, it simply gets lost in the machine (neither accelerated nor deflected).

The Cl+ ions will pass through the machine and will give lines at 35 and 37, depending on the isotope and you would get exactly the pattern in the last diagram. The problem is that you will also record lines for the unfragmented Cl2+ ions.

At the end the spectrum will show peaks due to ionized atoms, Cl+at 35, and 37, and ionized molecule Cl2+ at 70, 72, 74 as below

Checking up 1.5

The mass spectrum of magnesium is given below:

a. How many isotopes does magnesium possess

b. Estimate the isotopic mass of each of the magnesium isotopes

c. Estimate the relative abundance for each of the isotopes of magnesium

1.6. Uses of the mass spectrometer and involving calculations

Activity 1.6

1. Mass spectrometers are used to determine which of the following?

a) The atomic mass

b)Composition in sample

c) Concentration of elements in sample

2. The mass spectrum of an element, A, contained four lines at mass/charge of 54; 56; 57 and 58 with relative intensities of 5.84; 91.68; 542.17; 0.31 respectively. Explain these data and calculate the relative atomic mass of A

1.6.1. Calculation of RAM using mass spectrum

When the mass spectrum of the element is given, you can calculate the relative atomic mass of that element by using the information from the mass spectrum.

Example 1: the mass spectrum of boron is given below

Determine the relative atomic mass of boron

From the mass spectrum given, we have123 typical atoms of boron (sum of relative abundances). 23 of these would be 10B and 100 would be 11B.

The total mass of these would be (23 x 10) + (100 x 11) = 1330

The average mass of these 123 atoms would be 1330 / 123 = 10.8 (to 3 significant figures).10.8 is the relative atomic mass of boron.

Example 2: The figure below represents the mass spectrum of zirconium.

Suppose you had 100 typical atoms of zirconium (sum of relative abundances). 51.5 of these would be 90Zr, 11.2 would be 91Zr and so on. The total mass of these 100 typical atoms would be

(51.5 x 90) + (11.2 x 91) + (17.1 x 92) + (17.4 x 94) + (2.8 x 96) = 9131.8

The average mass of these 100 atoms would be 9131.8 / 100 = 91.3 (to 3 significant figures).

91.3 is the relative atomic mass of zirconium.

1.6.2. Uses of mass spectrometer

In addition to the use of mass spectrometer in the determination of isotopes of elements and their relative abundances, the applications of mass spectrometry are found:

•Pharmaceutical: drug discovery, combinatorial chemistry, pharmacokinetics, drug metabolism.

•Clinical: neonatal screening, haemoglobin analysis, drug testing.

•Environmental: water quality, soil and groundwater contamination, food contamination, pesticides on foods.

•Geological: oil composition.

•Biotechnology: the analysis of proteins, peptides.

Checking up 1.6

1. Which of the following is not done through mass spectrometry?

a. Calculating the isotopic abundance of elements

b. Investigating the elemental composition of planets

c. Confirming the presence of O-H and C=O in organic compounds

d. Calculating the molecular mass of organic compounds

2. Mass spectra enable you to find relative abundances of the isotopes of a particular element.

a) What are isotopes?

b) Define relative atomic mass.

c) The mass spectrum of strontium contains the following lines for 1+ ions:

Calculate the relative atomic mass of strontium.

3. The mass spectrum for chlorine looks like this:

a) Explain why there are two separate groups of peaks.

b) State what causes each of the 5 lines.

c) Explain the approximate relative heights of the lines at 35 and 37.

d) Why cannot you predict the relative heights of the two clusters of lines (35/37 and 70/72/74)?

1.7. End unit assessment

I. Multiple choice questions

1.Which of the following is true regarding a typical atom?

a. Neutrons and electrons have the same mass.

b. The mass of neutrons is much less than that of electrons.

c. Neutrons and protons together make the nucleus electrically neutral.

d. Protons are more massive than electrons

2. Which of the following statements is(are) true? For the false statements, correct them.

a. All particles in the nucleus of an atom are charged.

b. The atom is best described as a uniform sphere of matter in which electrons are embedded.

c. The mass of the nucleus is only a very small fraction of the mass of the entire atom.

d. The volume of the nucleus is only a very small fraction of the total volume of the atom.

e. The number of neutrons in a neutral atom must equal the number of electrons.

3. Each of the following statements is true, but Dalton might have had trouble

 explaining some of them with his atomic theory. Give explanations for the following statements.

a. Atoms can be broken down into smaller particles.

b. One sample of lithium hydride is 87.4% lithium by mass, while another 

sample of lithium hydride is 74.9% lithium by mass. However, 

the two samples have the same chemical properties

4. In mass spectrometer, the sample that has to be analysed is 

bombarded with which of the following?

a. Protons

b. Electrons

c. Neutrons

d. Alpha particles top of Form

5. Mass spectrometer separates ions on the basis of which of the following?

a. Mass

b. Charge

c. Molecular weight

d. Mass to charge ratio

6. In a mass spectrometer, the ions are sorted out in which of the following ways?

a. By accelerating them through electric field

b. By accelerating them through magnetic field

c. By accelerating them through electric and magnetic field

d. By applying a high voltage

7. The procedure for mass spectroscopy starts with which of the following processes?

a. The sample is bombarded by electron beam

c. The sample is converted into gaseous state

d. The ions are detected

8. Which of the following ions pass through the slit and reach the collecting plate?

a. Negative ions of all masses

b. Positive ions of all masses

c. Negative ions of specific mass

d. Positive ions of specific mass

9. Which of the following statements is not true about mass spectrometry?

a. Impurities of masses different from the one being analysed interferes with the result

b. It has great sensitivity

c. It is suitable for data storage

d. It is suitable for library retrieval

10. In a mass spectrometer, the sample gas is introduced into the highly 

evacuated spectrometer tube and it is ionised by the electron beam. a. True b. False.

Short and long answer questions

11. What are the three fundamental particles from which atoms are built? What are their 

electric charges? Which of these particles constitute the nucleus of an atom? 

Which is the least massive particle of the fundamental particles?

12. Verify that the atomic weight of lithium is 6.94, given the following information:

6Li, mass = 6.015121 u; percent abundance = 7.50%

7Li, mass = 7.016003 u; percent abundance = 92.50%13. T

he diagram below shows the main parts of a mass spectrometer.

a. Describe the different steps involved in taking a mass spectrum of a sample

b. (i) Which two properties of the ions determine how much they are deflected

 by the magnetic field? What effect does each of these properties have on the extent of deflection?

(ii) Of the three different ion streams in the diagram above, why is the ion stream C least deflected?

(iii) What would you have to do to focus the ion stream C on the detector?

c. Why is it important that there is a vacuum in the instrument?

d. Describe briefly how the detector works. Bottom of Form

14. A mass spectrum of a sample of indium shows two peaks at m/z = 113 and m/z = 115. 

The relative atomic mass of indium is 114.5. Calculate the relative abundances of these two isotopes. 

(b)The mass spectrum of the sample of magnesium contains three peaks with the mass-to-charge rations and relative intensities shown below

i. Explain why magnesium gives three peaks in mass spectrum?

ii. Use the information in the table above to calculate the accurate value for the relative atomic mass of magnesium

15. There exists 3 isotopes of oxygen that occur naturally with atomic mass 16, 17 and 18 with abundance 99.1% ; 0.89% and 0.01% respectively. Given that oxygen occurs naturally as diatomic molecule,

a. Predict the number of peaks that will be observed on the screen of mass spectrometer.

b. Show the molecular ions that are responsible of these peaks.